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Structure and Properties of Organic Molecules

Highland Hall Biochem Block Handout #2. Structure and Properties of Organic Molecules. Adapted from Chapter 2 of Organic Chemistry by L. G. Wade, Jr. Molecular Orbitals (MO’s). Definition: interaction between orbitals on different atoms

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Structure and Properties of Organic Molecules

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  1. Highland Hall Biochem Block Handout #2 Structure and Propertiesof Organic Molecules Adapted from Chapter 2 of Organic Chemistry by L. G. Wade, Jr.

  2. Molecular Orbitals (MO’s) • Definition: interaction between orbitals on different atoms • Generated by Linear Combination of Atomic Orbitals (LCAO) • Conservation of orbitals • # of MO’s created = # of AO’s combined • Interactions can be constructive or destructive. => Chapter 2

  3. Bonding MO : formed by constructive combination of AO’s with high e- density between the two nuclei. Results in MO with lower energy. Bonding MO (between two 1S orbitals) Chapter 2

  4. Antibonding MO: MO formed by destructive combination of AO’s with low e- density between the two nuclei Results in MOs higher in energy than the AO’s. Antibonding MO(between two 1S orbitals) Chapter 2

  5. Two atomic orbitals will always overlap so as to give 1 bonding MO and 1 antibonding MO. Example: H2 MO Energy Level Diagram => Chapter 2

  6. 1S 1S The σ bond • Sigma (σ) bond : single bond formed by constructive overlap of orbitals in which the e- density is centered about a line connecting the nuclei (i.e. cylindrically symmetrical) • Constructive orbital overlap → σ bonding MO • Destructive orbital overlap → σ* anti-bonding MO Chapter 2

  7. π Bonding Pi (π) bond : formed by constructive side-to-side overlap of parallel p orbitals. Chapter 2 7

  8. π Bonding • Not cylindrically symmetrical • Parallel overlap required for π bonds • π bonds cannot rotate with respect to each other • π bonds formed in multiple bonds • π bonds form after a σ bond. Chapter 2

  9. => Multiple Bonds • A double bond (2 pairs of shared electrons) consists of a sigma bond and a pi bond. • A triple bond (3 pairs of shared electrons) consists of a sigma bond and two pi bonds. Chapter 2

  10. σ bonds vs π bonds Chapter 2

  11. Hybrid Orbitals • Hybrid orbitals : formed by mixing orbitals on the same atom. • Mixtures of s and p orbitals • Why do we need hybrids? • Bond angles cannot be explained with simple s and p orbitals, if so we would predict carbon to have 90o angles because 2px, 2py, and 2pz are orthogonal. • Actual angles are ~ 109o, 120o or 180o

  12. VSEPR • Valence Shell Electron Pair Repulsion (VSEPR) theory : e- pairs will repel each so as to attain a maximal distance of separation. • The observed bond angles are not possible using only s & p orbitals thus we use hybrids. • Hybridized orbitals are lower in energy because electron pairs are farther apart. • Hybridization is LCAO within one atom, just prior to bonding. Chapter 2

  13. Hybridization Rules 1) Both σ bonding e- and lone pairs occupy hybrid orbitals: # hybrid orbitals = (# σ bonds) + (# lone pairs) 2) The observed hybridization will be that which gives maximum separation of σ bonds and lone pairs 3) For multiple bonds, one bond is a σ bond using hybrid orbitals, while additional bonds are π bonds formed between p orbitals. Chapter 2

  14. sp Hybrid Orbitals(add an s and p orbital together) • 2 atomic orbitals = 2 hybrid orbitals • Linear electron pair geometry • 180° bond angle Chapter 2

  15. sp Hybrid Orbitals(add an s and p orbital together) Example: BeH2 Bond angle is 180o Chapter 2 15

  16. sp2 Hybrid Orbitals(add an s and 2 p orbitals together) • 3 atomic orbitals = 3 hybrid orbitals • Trigonal planar e- pair geometry • 120° bond angle Chapter 2

  17. sp2 Hybrid Orbitals(add an s and 2 p orbitals together) • Example: BeH3 • Bond angle is 120o Chapter 2 17

  18. sp3 Hybrid Orbitals(add an s and 3 p orbitals together) • 4 atomic orbitals = 4 hybrid orbitals • Tetrahedral e- pair geometry • 109.5° bond angle • Example : methane (CH4) Chapter 2

  19. Summary of s-p type orbitals

  20. Sample Problems • Predict the hybridization, geometry, and bond angle for each atom in the following molecules: • Caution! You must start with a good Lewis structure! NH2NH2 CH3-CC-CHO Chapter 2

  21. Isomers • Isomers : molecules which have the same molecular formula, but differ in the arrangement of their atoms • Constitutional (or structural) : isomers which differ in their bonding sequence. • Stereoisomers : molecules that differ only in the arrangement of the atoms in space. Chapter 2

  22. Structural Isomers Chapter 2

  23. Conformations • Conformations : structures differing only in rotations about a single bond. • Example: ethane • Cylindrical symmetry of the σ bond allows rotation about a single bond “eclipsed” “staggered” side view side view

  24. Rotation around Bonds • σ bonds freely rotate. • π bonds cannot rotate unless the bond is broken. • only one conformation • leads to stereoisomers Chapter 2

  25. Stereoisomers Trans - across Cis - same side Stereoisomer : distinct compounds differing only in the spatial arrangement of atoms attached to the C=C double bond. Cis-trans isomers are also called geometric isomers. There must be two different groups on the sp2 carbon. No cis-trans isomers possible Chapter 2 25

  26. Dipole Moment (μ) • Polar bond : bond in which e- are shared unequally. • Non-polar : bond in which e- are shared equally. • Dipole moment (μ) : measure of polarity. • are due to differences in electronegativity. • depend on the amount of charge and distance of separation. • In debyes (D),  = 4.8 x (electron charge) x d(angstroms)

  27. Bond Dipole Moments(μ)

  28. Molecular Dipole Moments • Depends on bond polarity and bond angles. • Vector sum of the bond dipole moments. • Lone pairs of electrons contribute to the dipole moment. μ = 1.9 D μ = 2.3 Dμ = 0 D Chapter 2

  29. More examples: Chapter 2 29

  30. Intermolecular Forces • Strength of attractions between molecules influence m.p., b.p., and solubility; especially for solids and liquids. • Classification depends on structure. • Dipole-dipole interactions • London dispersions • Hydrogen bonding => Chapter 2

  31. Dipole-Dipole Forces • Between polar molecules • Positive end of one molecule aligns with negative end of another molecule. • Lower energy than repulsions, so net force is attractive. • Larger dipoles cause higher boiling points and higher heats of vaporization. Chapter 2

  32. Dipole-Dipole => Chapter 2

  33. London Dispersion • Induced between non-polar molecules • Temporary dipole-dipole interactions • Forces are larger for molecules with large surface area. • large molecules are more polarizable • Branching lowers b.p. because of decreased surface area contact between molecules. Chapter 2

  34. Dispersions => Chapter 2

  35. Hydrogen Bonding Strong dipole-dipole attraction Organic molecule must have N-H or O-H. The hydrogen from one molecule is strongly attracted to a lone pair of electrons on the other molecule. O-H more polar than N-H, so there is stronger hydrogen bonding. Chapter 2 35

  36. H Bonds => Chapter 2

  37. Hydrogen Bonding & BP Higher bp’s for NH3 & H2O which can H-bond H2O bp is so high because… Chapter 2 37

  38. => Boiling Points and Intermolecular Forces Chapter 2

  39. Solubility • The solubility rule: “Like dissolves like” • Ionic & polar solutes dissolve in polar solvents. • Nonpolar solutes dissolve in nonpolar solvents. • Molecules with similar intermolecular forces will mix freely. Chapter 2

  40. Ionic Solute with Polar Solvent Hydration releases energy. Entropy increases. => Chapter 2

  41. Ionic Solute withNonpolar Solvent Attraction for solvent much weaker than ion-ion interactions Chapter 2

  42. Nonpolar Solute withNonpolar Solvent Small attractive forces for solvent are still strong enough to over come weak ion-ion intermolecular forces in solid. Chapter 2

  43. Nonpolar Solute with Polar Solvent Dissolving solute would require breaking up organized solvent structure. (crisco in water) Chapter 2

  44. Classes of Compounds • Classification based on functional group • Three broad classes • Hydrocarbons • Compounds containing oxygen • Compounds containing nitrogen => Chapter 2

  45. Hydrocarbons • Compounds containing only C and H. • Includes alkane, cycloalkane, alkene, cycloalkene, alkyne, and aromatic hydrocarbons • Typically hydrophobic Chapter 2

  46. Alkanes Hydrocarbons that contain only single bonds, sp3 carbons (CH3-CH2-CH3) Can be cyclic : then called cycloalkane Alkanes are very unreactive Alkane portion of molecule is called an alkyl group Chapter 2 47

  47. Alkenes Hydrocarbons that contain C=C double bonds, sp2 carbons. More reactive than alkanes Show geometric isomerism Cycloalkene: double bond in ring Chapter 2 48

  48. Alkynes Hydrocarbons with C—C triple bond, sp carbons More reactive than alkanes & alkenes Chapter 2 49

  49. Aromatic Hydrocarbons Hydrocarbons with alternating C—C single bond and C=C double bonds. Have unusual stability compared to alkenes and cycloalkenes Chapter 2 50

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